Animated Solution for Chemistry - Ionic Equilibrium: What is the conjugate base of ${\text{OH}}^{-}$?

Have you ever wondered how molecules and ions transform into one another just by passing around a single, tiny particle? In the world of acid-base chemistry, the proton (a hydrogen ion, ${\text{H}}^{+}$) is the ultimate currency. It is traded back and forth, and with every transaction, the identity of the molecule changes.

Today, we are going to dive into a classic, fundamental question: What is the conjugate base of ${\text{OH}}^{-}$?

While it might look like a simple one-step problem, the concept behind it—the Brønsted-Lowry theory—is a cornerstone of chemistry. Let's break it down step by step and truly understand the mechanics of proton transfer.

To find a conjugate base, we first need to know what it means. In 1923, Johannes Nicolaus Brønsted and Thomas Martin Lowry independently proposed a revolutionary definition for acids and bases.

They stated that an acid is any species that can donate a proton (${\text{H}}^{+}$), and a base is any species that can accept a proton.

When an acid donates its proton, it doesn't just disappear into the void. The acid transforms into a new species. Because this new species now has the potential to accept a proton back (reversing the reaction), it acts as a base. We call this the conjugate base of the original acid.

Mathematically, we can write this as:

In our specific problem, we are given the hydroxide ion, ${\text{OH}}^{-}$, and asked to find its conjugate base.

This phrasing is a direct instruction: "Treat ${\text{OH}}^{-}$ as an acid."

Wait, isn't ${\text{OH}}^{-}$ the classic example of a base? Yes, in Arrhenius chemistry, it is the ultimate base! But in the Brønsted-Lowry world, many species are amphoteric, meaning they can act as either an acid or a base depending on the situation. Water (${\text{H}}_2\text{O}$) is the most famous amphoteric species, but ${\text{OH}}^{-}$ can also play both roles.

To find the conjugate base, we must force ${\text{OH}}^{-}$ to act as an acid. It must donate a proton.

Let's set up our chemical equation. We start with our reactant:

Now, we extract one proton (${\text{H}}^{+}$) from it.

What is that "Something"? We need to balance two things: the atoms and the electrical charge.

1. Balancing the Atoms: The hydroxide ion has one Oxygen atom and one Hydrogen atom. If we remove one Hydrogen atom (as ${\text{H}}^{+}$), we are left with only the Oxygen atom.

2. Balancing the Charge: This is where silly mistakes often happen. The original ${\text{OH}}^{-}$ ion has a net charge of $-1$. We are removing a proton, which carries a charge of $+1$. Mathematically, subtracting a positive charge makes the remaining species more negative:

Therefore, the Oxygen atom must carry a charge of $-2$.

Putting the atoms and the charge together, the species left behind is the oxide ion:

Let's write the complete, balanced equation for this proton donation:

The oxide ion, ${\text{O}}^{2-}$, is the conjugate base of the hydroxide ion. Looking at our options, this matches perfectly with option (a).

Conjugate Acids

To solidify this concept, let's flip the script. What if the question had asked for the conjugate acid of ${\text{OH}}^{-}$?

To find a conjugate acid, we must treat the species as a base. It must accept a proton.

Adding one Hydrogen atom to $\text{OH}$ gives us ${\text{H}}_2\text{O}$. Adding a $+1$ charge to a $-1$ charge gives a net charge of $0$. So, the conjugate acid of ${\text{OH}}^{-}$ is neutral water, ${\text{H}}_2\text{O}$.

The Golden Rule: A conjugate acid-base pair always differs by exactly one ${\text{H}}^{+}$ ion. - To find the conjugate base: Subtract ${\text{H}}^{+}$ - To find the conjugate acid: Add ${\text{H}}^{+}$

By mastering this simple accounting of atoms and charges, you can confidently navigate any Brønsted-Lowry acid-base question that comes your way!